Ions in Solids and Solutions

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Chemistry F332 Notes
Ions in solids and solutions:
Structure of an ionic lattice (Sodium Chloride):
* Consists of sodium ions (Na+) surrounded by six chloride ions (Cl-) * Chloride ions also surrounded by six sodium ions.
* Held together by attraction of oppositely charged ions. * Giant ionic lattice.
* Electrostatic bonds hold lattices together.
* Structure is simple cubic.
* Some ionic crystals contain water.
* Known as water of crystallisation.
* These crystals are hydrated.
* Blue copper sulphate crystals are the pentahydrate crystals – CuSO4.5H2O * If there is no water of crystallisation then the ionic crystal is said to be anhydrous. Ionic substances in solution:

* Many dissolve in water without difficulty.
* Ions become surrounded by polar water molecules.
* They spread out through the solution.
* Hydrated ions are randomly arranged and behave independently. * Positive hydrogen atoms in water are attracted to negative ions. * Conversely the negative oxygen atoms are attracted to positive ions. * This is known as hydration.

* Draw water molecules as wedges and ion as a circle.
* Label the charged ends.
Ionic equations:
* All nitrates are soluble in water.
* All chlorides are soluble in water bar AgCl and PbCl2.
* All sulfates are soluble in water bar BaSO4, PbSO4 and SrSO4. * All sodium, potassium and ammonium salts are soluble in water. * All carbonates are insoluble in water bar (NH4)2CO3 and those of group 1. * The ionic equation for a neutralisation reaction is always – H+(aq) + OH-(aq) → H2O(l) * When writing an ionic equation label the charges on each ion then remove the spectator ions (like ions in charge and state and on both sides of equation). * If there is a precipitation reaction one product will be a solid. * With neutralisation reacts a covalent compound is formed. Atoms and ions:

First ionisation enthalpy:
* Energy needed to remove one electron from each of one mole of isolated gaseous atoms of an element. * One mole of gaseous ions with one positive charge is formed: X(g) →X+(g) + e- * The value of this is always positive.

* Energy must be put in to remove electron which is attracted to the nucleus. * Largest enthalpies are noble gases, they have a full outer shell and all very unreactive. * Lowest enthalpies for group 1 elements, very reactive elements and only have 1 electron in there outer shell. * Across a period enthalpies increase as an extra electron is being added into the same shell, therefore the nuclear charge is increasing. Thus a greater attraction attraction the nucleus and electron. Successive ionisation enthalpies:

* Ionisation enthalpies which are the energies required to remove further electrons * Second: X+(g) → X2+(g) + e-
* Enthalpies increases as successive electrons are removed * After an electron is removed, the remaining electrons are held on more firmly * There is a sharp jump when an electron is removed from a full electron shell. Oxidation number rules:

* All elements have an oxidation number of 0.
* For simple ions the oxidation number is the charge the ion carries. * In a neutral compound the sum of the oxidation numbers is 0. * (K+)2O2- → (+1)2 -2 = 0
* Fluorine (the most electronegative element) always has an oxidation number of -1. * Oxygen always has a oxidation number of -2 except…
* F2O →(-1)2 +2 = 0
* Na2O2 → (+1)2 (-1)2 = 0
* Hydrogen always has an oxidation number of +1, except in metallic hydrides when it is -1. * In complex ions the sum of the oxidation numbers = the charge the ion carries. Oxyacids –

H2SO4 → Sulphuric(vi) acid.
H2SO3 → Sulphuric(iv) acid.
HNO3 → Nitric(v) acid.
HClO → Chloric(i) acid.
HClO3 → Chloric(v) acid.
HClO4 → Chloric(vii) acid.

Oxyacids –

H2SO4 → Sulphuric(vi) acid.
H2SO3 → Sulphuric(iv) acid.
HNO3 → Nitric(v) acid.
HClO → Chloric(i)...
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